CHEMISTRY

1. INTRODUCTION

The syllabus is designed for examination purposes, so the order of topics is not necessarily the order of teaching.
Assumptions:
- Candidates have covered Integrated Science/Basic Science or General Science and Mathematics at the Junior Secondary School (JSS)/Junior High School (J.H.S) level.
- Candidates have engaged in suggested activities and project work to develop competencies and skills outlined in Chemistry teaching syllabuses.
- Schools offering the subject have well-equipped laboratories.
Candidates should know significant figures, S.I. units, and the conventional/IUPAC system of nomenclature.

2. AIMS

The syllabus aims to assess:

Understanding of basic chemistry concepts.
Acquisition of laboratory skills, including awareness of hazards and safety measures.
Awareness of the inter-relationship between chemistry and other disciplines.
Awareness of the linkage between chemistry and industry/environment/everyday life in terms of benefits and hazards.
Skills of critical and logical thinking.

3. EXAMINATION SCHEME

Three papers: Papers 1, 2, and 3.
Papers 1 and 2 are a composite paper taken at one sitting.

PAPER 1:

Fifty multiple-choice objective questions from Section A of the syllabus (common to all candidates).
1 hour, 50 marks.

PAPER 2:

2-hour essay paper covering the entire syllabus, 100 marks.
Two sections: A and B.
- Section A: Ten short structured questions from the common portion of the syllabus (Section A), 25 marks.
- Section B: Two questions from the common portion (Section A) and two from the country-specific section (Section B or C). Candidates answer any three questions, each worth 25 marks.

PAPER 3:

2-hour practical test for school candidates or 1 hour 30 minutes alternative to practical work test for private candidates, 50 marks.
Three compulsory questions:
- One on quantitative analysis.
- One on qualitative analysis.
- One testing familiarity with practical activities suggested in teaching syllabuses.
Details of continuous assessment input given by the Council.

SECTION A (For all candidates)

1.0 INTRODUCTION TO CHEMISTRY

(a) (i) Measurement of physical quantities:
- Measurement of mass, length, time, temperature, and volume.
- Appropriate SI units and significant figures.
- Precision and accuracy in measurement.
(b) Scientific Methods:
- Outline the scientific method: Observation, hypothesis, experimentation, formulation of laws, and theories.

2.0 STRUCTURE OF THE ATOM

(a) Gross features of the atom:
- Short account of Dalton’s atomic theory and its limitations, J.J. Thompson’s experiment, and Bohr’s model of the atom.
- Outline description of Rutherford’s alpha scattering experiment to establish the structure of the atom.
(b) (i) Atomic number/proton number, number of neutrons, isotopes, atomic mass, mass number:
- Meaning and representation in symbols of atoms and sub-atomic particles.
(ii) Relative atomic mass (Ar) and relative molecular mass (Mr) based on Carbon-12 scale:
- Atomic mass as the weighted average mass of isotopes. Calculation of relative mass of chlorine should be used as an example.
- Carbon-12 scale as a unit of measurement.
- Definition of atomic mass unit.
(iii) Characteristics and nature of matter:
- Atoms, molecules, and ions.
(c) Particulate nature of matter: physical and chemical changes:
- Definition of particles and treatment of particles as building blocks of matter.
- Explain physical and chemical changes with examples.
  - Physical change: melting of solids, magnetization of iron, dissolution of salt, etc.
  - Chemical change: burning of wood, rusting of iron, decay of leaves, etc.
(d) (i) Electron Configuration:
- Detailed electron configurations (s, p, d) for atoms of the first thirty elements.
(ii) Orbitals:
- Origin of s, p, and d orbitals as sub-energy levels; shapes of s and p orbitals only.
(iii) Rules and principles for filling in electrons:
- Aufbau Principle, Hund’s Rule of Maximum Multiplicity, and Pauli Exclusion Principle.
- Abbreviated and detailed electron configuration in terms of s, p, and d.

3.0 STANDARD SEPARATION TECHNIQUES FOR MIXTURES

(a) Classification of mixtures:
- Solid-solid, solid-liquid, liquid-liquid, gas-gas with examples.
(b) Separation techniques:
- Crystallization, distillation, precipitation, magnetization, chromatography, sublimation, etc.
(c) Criteria for purity:
- Boiling point for liquids and melting point for solids.

4.0 PERIODIC CHEMISTRY

(a) Periodicity of the elements:
- Electron configurations leading to group and periodic classifications.
(b) Different categories of elements in the periodic table:
- Metals, semi-metals, non-metals in the periodic table and halogens. Alkali metals, alkaline earth metals, and transition metals as metals.
(c) Periodic law:
- Explanation of the periodic law.
- (i) Trends on periodic table:
  - Periodic properties: atomic size, ionic size, ionization energy, electron affinity, and electronegativity.
  - Simple discrepancies should be accounted for in respect to beryllium, boron, oxygen, and nitrogen.
- (ii) Periodic gradation of the elements in the third period (Na - Ar):
  - Progression from:
    - metallic to non-metallic character of element;
    - ionic to covalent bonding in compounds.
  - Differences and similarities in the properties between the second and the third-period elements should be stated.
(d) Reactions between acids and metals, their oxides and trioxocarbonates (IV):
- Period three metals (Na, Mg, Al).
- Period four metals (K, Ca).
- Chemical equations.
- pH of solutions of the metallic oxides and trioxocarbonates.
(e) Periodic gradation of elements in group seven, the halogens: F, Cl, Br, and I:
- Recognition of group variations noting any anomalies. Treatment should include the following:
  - physical states, melting and boiling points;
  - variable oxidation states;
  - redox properties of the elements;
  - displacement reaction of one halogen by another;
  - reaction of the elements with water and alkali (balanced equations required).
(f) Elements of the first transition series. 21Sc – 30Zn:
- Their electron configurations, physical properties, and chemical reactivity of the elements and their compounds.
- Physical properties should include: physical states, metallic properties, and magnetic properties.
- Reactivity of the metals with air, water, acids, and comparison with s-block elements (Li, Na, Be, Mg).
- Other properties of transition metals should include:
  - variable oxidation states;
  - formation of colored compounds;
  - complex formation;
  - catalytic abilities;
  - paramagnetism;
  - hardness.

5.0 CHEMICAL BONDS

(a) Interatomic bonding:
- Meaning of chemical bonding. Lewis dot structure for simple ionic and covalent compounds.
(b) (i) Formation of ionic bonds and compounds:
- Formation of stable compounds from ions. Factors influencing formation: ionization energy; electron affinity and electronegativity difference.
- (ii) Properties of ionic compounds:
  - Solubility in polar and non-polar solvents, electrical conductivity, hardness, and melting point.
(c) Naming of ionic compounds:
- IUPAC system for simple ionic compounds.
(d) Formation of covalent bonds and compounds:
- Factors influencing covalent bond formation. Electron affinity, ionization energy, atomic size, and electronegativity.
(e) (i) Properties of covalent compounds:
- Solubility in polar and non-polar solvents, melting point, boiling point, and electrical conductivity.
- (ii) Coordinate (dative) covalent bonding:
  - Formation and difference between pure covalent and coordinate (dative) covalent bonds.
(f) Shapes of molecular compounds:
- Linear, planar, tetrahedral and shapes for some compounds e.g. BeCl2, BF3, CH4, NH3, CO_2. Factors should include: atomic radius, ionization energy, and number of valence electrons.
(g) (i) Metallic Bonding:
- Factors influencing its formation. Factors should include: atomic radius, ionization energy, and number of valence electrons. Types of specific packing not required.
  - (ii) Properties of metals:
    - Typical properties, including heat and electrical conductivity, malleability, luster, ductility, sonority, and hardness.
(h) (i) Intermolecular bonding:
- Relative physical properties of polar and non-polar compounds.
  - (ii) Intermolecular forces in covalent compounds:
    - Description of formation and nature should be treated.
    - (iii) Hydrogen bonding:
      - (iii) van der Waals forces:
        Dipole-dipole, induced dipole-dipole, induced dipole-induced dipole forces should be treated under van der Waal’s forces.
        Variation of the melting points and boiling points of noble gases, halogens, and alkanes in the homologous series explained in terms of van der Waal’s forces; and variation in the boiling points of H2O and H2S explained using Hydrogen bonding.
- **(iv) Comparison of all bond types.

6.0 STOICHIOMETRY AND CHEMICAL REACTIONS

(a) (i) Symbols, formulae, and equations:
- Symbols of the first thirty elements and other common elements that are not among the first thirty elements.
**(ii) chemical symbols
(iii) Empirical and molecular formulae:
- Calculations involving formulae and equations will be required. Mass and volume relationships in chemical reactions and the stoichiometry of such reactions such as: calculation of percentage composition of element.
(iv) Chemical equations and IUPAC names of chemical compounds:
- Combustion reactions (including combustion of simple hydrocarbons)
- Synthesis
- Displacement or replacement
- Decomposition
- Ionic reactions
(v) Laws of chemical combination:
- Laws of conservation of mass.
- Law of constant composition.
- Law of multiple proportions. Explanation of the laws to balance given equations.
- Experimental illustration of the law of conservation of mass.
(b) Amount of substance:
- Mass and volume measurements.
- The mole as a unit of measurement; Avogadro’s constant, L = 6.02 \times 10^{23} entities \; mol^{-1}.
- Molar quantities and their uses.
- Moles of electrons, atoms, molecules, formula units, etc.
(c) Mole ratios:
- Use of mole ratios in determining stoichiometry of chemical reactions.
- Simple calculations to determine the number of entities, amount of substance, mass, concentration, volume, and percentage yield of product.
(d) (i) Solutions:
- Concept of a solution as made up of solvent and solute.
- Distinguishing between dilute solution and concentrated solution.
- Basic, acidic, and neutral solutions.
- (ii) Concentration terms:
  - Mass (g) or moles (mol) per unit volume. Emphasis on current IUPAC chemical terminology, symbols, and conventions. Concentration be expressed as mass concentration, g dm-3, molar concentration, mol dm-3.
  - (iii) Standard solutions:
    - Preparation of some primary standards e.g., anhydrous Na2CO3, (COOH)2.2H2O / H2C2O4.2H2O.
    - Meaning of the terms primary standard, secondary standard, and standard solution.
(e) Preparation of solutions from liquid solutes by the method of dilution:
- Dilution factor.

7.0 STATES OF MATTER

(a) (i) Kinetic theory of matter:
- Postulates of the kinetic theory of matter.
- Use of the kinetic theory to explain the following processes: melting of solids, boiling of liquids, evaporation of liquids, dissolution of solutes, Brownian motion, and diffusion.
(ii) Changes of state of matter:
- Changes of state of matter should be explained in terms of movement of particles. It should be emphasized that randomness decreases (and orderliness increases) from the gaseous state to the liquid state and to the solid state and vice versa.
- Illustrations of changes of state using the different forms of water, iodine, sulfur, naphthalene, etc.
- Brownian motion to be illustrated using any of the following experiments:
  - pollen grains/powdered sulfur in water (viewed under a microscope);
  - smoke in a glass container illuminated by a strong light from the side;
  - a dusty room being swept and viewed from outside under sunlight.
(iii) Diffusion:
- Experimental demonstration of diffusion of two gases.
- Relationship between the speed at which different gas particles move and the masses of particles.
- Experimental demonstration of diffusion of solute particles in liquids.
(b) Gases:
- (i) Characteristics and nature of gases:
  - Arrangement of particles, density, shape, and compressibility.
- (ii) The gas laws:
  - Charles’ Law, Boyle’s Law, Dalton’s law of partial pressure, Graham’s law of diffusion, Avogadro’s law. The ideal gas equation of state.
  - Qualitative explanation of each of the gas laws using the kinetic model. The use of Kinetic molecular theory to explain changes in gas volumes, pressure, temperature.
  - Mathematical relations of the gas law PV= nRT. Ideal and Real gases. Factors responsible for the deviation of real gases from the ideal situation.
- (iii) Laboratory preparation and properties of some gases:
  - Preparation of the following gases: H2, NH3, and CO_2. Principles of purification and collection of gases.
  - Physical and chemical properties of the gases.
(c) (i) Liquids:
- Characteristics and nature of liquids based on the arrangement of particles, shape, volume, compressibility, density, and viscosity. (ii) Vapour and gases:
  - Concept of vapor, vapor pressure, saturated vapor pressure, boiling, and evaporation.
  - Distinction between vapor and gas.
  - Effect of vapor pressure on boiling points of liquids.
  - Boiling at reduced pressure.
(d) Solids:
- **(i) Characteristics and nature:
- (ii) Types and structures:
  - Ionic, metallic, covalent network, and molecular solids. Examples in each case.
  - Arrangements of particles ions, molecules and atoms in the solid state.
- (iii) Properties of solids:
  - Relate the properties of solids to the type of interatomic and intermolecular bonding in the solids. Identification of the types of chemical bonds in graphite and differences in the physical properties.
(e) Structures, properties and uses of diamond and graphite:
- The uses of diamond and graphite related to the structure.
- The use of iodine in everyday life.
(f) Determination of melting points of covalent solids:
- Melting points as an indicator of purity of solids e.g., Phenyl methanedioic acid (benzoic acid), ethanedioic acid (oxalic), and ethanamide.

8.0 ENERGY AND ENERGY CHANGES

(a) Energy and enthalpy:
- Explanation of the terms energy and enthalpy. Energy changes associated with chemical processes.
(b) Description, definition, and illustrations of energy changes and their effects:
- Exothermic and endothermic processes.
- Total energy of a system as the sum of various forms of energy e.g., kinetic, potential, electrical, heat, sound, etc.
- Enthalpy changes involved in the following processes: combustion, dissolution, and neutralization.

9.0 ACIDS, BASES, AND SALTS

(a) Definitions of acids and bases:
- Arrhenius concepts of acids and bases in terms of H_3O^+ and OH^- ions in water.
- Effects of acids and bases on indicators, metal Zn, Fe, and trioxocarbonate (IV) salts and hydrogentrioxocarbonate (IV) salts.
(b) Physical and chemical properties of acids and bases:
- Characteristic properties of acids and bases in aqueous solution to include:
  - conductivities, taste, litmus/indicators, feel etc.;
  - balanced chemical equations of all reactions.
(c) Acids, bases, and salts as electrolytes:
- Electrolytes and non-electrolytes; strong and weak electrolytes. Evidence from conductivity and enthalpy of neutralization.
(d) Classification of acids and bases:
- Strength of acids and bases.
- Classify acids and bases into strong and weak.
- Extent of dissociation reaction with water and conductivity.
- Behavior of weak acids and weak bases in water as an example of equilibrium systems.
(e) Concept of pH:
- Definition of pH and knowledge of the pH scale.
- Measurement of pH of solutions using a pH meter, colorimetric methods, or universal indicator.
- Significance of pH values in everyday life e.g., acid rain, pH of soil, blood, urine.
(f) Salts:
- Meaning of salts.
- Types of salts: normal, acidic, basic, double, and complex salts.
  - (i) Laboratory and industrial preparation of salts:
    - Description of laboratory and industrial production of salts.
    - Mining of impure sodium chloride and conversion into granulated salt.
    - Preparation of NaOH, Cl2, and H2.
  - **(ii) Uses
  - (iii) Hydrolysis of salt:**
    - Explanation of how salts form acidic, alkaline, and neutral aqueous solutions.
    - Behavior of some salts (e.g., NH4Cl, AlCl3, Na2CO3, CH_3COONa) in water as examples of equilibrium systems.
    - Effects of charge density of some cations and anions on the hydrolysis of their aqueous solution. Examples to be taken from group 1, group 2, group 3, and the d-block element.
(g) Deliquescent, efflorescent, and hygroscopic compound:
- Use of hygroscopic compounds as drying agents should be emphasized.
(h) Acid-Base indicators:
- Qualitative description of how acid-base indicator works.
- Indicators as weak organic acids or bases (organic dyes).
- The color of the indicator at any pH depends on the relative amounts of acid and forms.
- Working pH ranges of methyl orange and phenolphthalein.
(i) Acid-Base titration:
- Knowledge and correct use of relevant apparatus.
- Knowledge of how acid-bases indicators work in titrations.
- Acid-base titration experiments involving HCl, HNO3, H2SO4, and NaOH, KOH, Ca(OH)2, CO3^{2-}, HCO3^-.
Titration involving weak acids versus strong bases, strong acids versus weak bases and strong acids versus strong bases using the appropriate indicators and their applications in quantitative determination; e.g. concentrations, mole ratio, purity, water of crystallization and composition.

10.0 SOLUBILITY OF SUBSTANCES

(a) General principles:
- Meaning of Solubility.
- Saturated and unsaturated solutions.
- Saturated solution as an equilibrium system.
- Solubility expressed in terms of: mol dm-3 and g dm-3 of solution/solvent.
- Solubility curves and their uses.
- Effect of temperature on solubility of a substance.
- Relationship between solubility and crystallization.
- Crystallization/recrystallization as a method of purification.
- Knowledge of soluble and insoluble salts of stated cations and anions.
- Calculations on solubility.
(b) Practical application of solubility:
- Generalization about the solubility of salts and their applications to qualitative analysis e.g., Pb^{2+}, Ca^{2+}, Al^{3+}, Cu^{2+}, Fe^{2+}, Fe^{3+}, Cl^{-}, Br^{-}, I^{-}, SO4^{2-}, S^{2-}, CO3^{2-}, Zn^{2+}, NH4^+, SO3^{2-}. Explanation of solubility rules.

11.0 CHEMICAL KINETICS AND EQUILIBRIUM SYSTEM

(a) Rate of reactions:
- Definition of reaction rate.
- Observable physical and chemical changes: color, mass, temperature, pH, the formation of precipitate etc.
  - (i) Factors affecting rates:
    - Physical states, concentration/ pressure of reactants, temperature, catalysts, light, particle size and nature of reactants.
    - Appropriate experimental demonstration for each factor is required.
  - (ii) Theories of reaction rates:
    - Collision and transition state theories to be treated qualitatively only.
    - Factors influencing collisions: temperature and concentration.
    - Effective collision.
    - Activation energy.
    - Energy profile showing activation energy and enthalpy change.
  - (iii) Analysis and interpretation of graphs:
    - Drawing of graphs and charts.
(b) Equilibrium:
- Explanation of reversible and irreversible reactions. Reversible reaction i.e., dynamic equilibrium.
- Equilibrium constant K must be treated qualitatively. It must be stressed that K for a system is constant at constant temperature. Simple experiment to demonstrate reversible reactions.
  - (i) General Principle:
    - Prediction of the effects of external influence of concentration, temperature pressure and volume changes on equilibrium systems.
      Le Chatelier’s principle.

12.0 REDOX REACTIONS

(a) Oxidation and reduction process:
- Oxidation and reduction in terms of:
  - addition and removal of oxygen and hydrogen;
  - loss and gain of electrons;
  - change in oxidation numbers/states.
- Determination of oxidation numbers/states.
(b) Oxidizing and reducing agents:
- Description of oxidizing and reducing agents in terms of:
  - addition and removal of oxygen and hydrogen;
  - loss and gain of electrons;
  - change in oxidation numbers/state.
(c) Redox equations:
- Balancing redox equations by:
  - ion, electron, or change in oxidation number/states;
  - half reactions and overall reaction.
(d) Electrochemical cells:
- Definition/Explanation
  - (i) Standard electrode potential:
    - Standard hydrogen electrode: meaning of standard electrode potential (Eo) and its measurement.
    - Only metal/metal ion systems should be used.
  - (ii) Drawing of cell diagram and writing cell notation:
  - (iii) e.m.f of cells:
    - Electrochemical cells as a combination of two half-cells.
    - The meaning of magnitude and sign of the e.m.f.
  - (iv) Application of Electrochemical cells:
    - Distinction between primary and secondary cells
    - Daniell cell, lead acid battery cell, dry cells, fuel cells and their use as generators of electrical energy from chemical reactions.
(e) Electrolysis:
- Definition.
  - (i) Electrolytic cells:
    - Comparison of electrolytic and electrochemical cells; weak and strong electrolyte.
  - (ii) Principles of electrolysis:
    - Mechanism of electrolysis.
  - (iii) Factors influencing the discharge of species:
    - Limit electrolytes to molten PbBr2 and NaCl, dilute NaCl solution, concentrated NaCl solution, CuSO4(aq), dilute H2SO4, NaOH(aq), and CaCl_2(aq) (using platinum or graphite and copper electrodes).
  - (iv) Faraday’s laws:
    - Simple calculations based on the relation 1F= 96,500 C and mole ratios to determine mass, volume of gases, number of entities, charges etc. using half and overall reactions.
  - (v) Practical application:
    - Electroplating, extraction, and purification of metals.
  - (vi) Corrosion of metals:
    - Corrosion treated as a redox process.
    - Rusting of iron and its economic costs.
    - Prevention based on the relative magnitude of electrode potentials and preventive methods like galvanizing, sacrificial/cathodic protection, and non-redox methods (painting, greasing/oiling etc.).

13.0 CHEMISTRY OF CARBON COMPOUNDS

(a) Classification:
- Broad classification into straight-chain, branched-chain, aromatic, and alicyclic compounds.
(b) Functional group:
- Systematic nomenclature of compounds with the following functional groups: alkanes, alkenes, alkynes, hydroxyl compounds (aliphatic and aromatic), alkanoic acids, alkyl alkanoates (esters and salts), and amines.
(b) Separation and purification of organic compounds:
- Methods to be discussed should include: distillation; crystallization; drying and chromatography.
(c) Petroleum/crude oil:
- Composition and classification.
- Fractional distillation and major products.
- Cracking and reforming.
- Petro-chemicals: sources; uses e.g., as starting materials of organic synthesis.
- Quality of petrol, meaning of octane number, and its importance to the petroleum industry.
**(d) Determination of empirical and molecular formulae and molecular structures of organic compounds.
(e) General properties of organic compounds:
- (i) Homologous series:
  - Gradation in physical properties.
- (ii) Isomerism:
  - Effects on the physical properties by introduction of active groups into the inert alkane.
  - Examples should be limited to compounds having a maximum of five carbon atoms.
  - Differences between structural and geometric/stereo isomerism.
(f) Alkanes:
- (i) Sources, properties:
  - Laboratory and industrial preparations and other sources.
  - Nomenclature and structure.
  - Reactivity:
    - combustion;
    - substitution reactions;
    - cracking of large alkane molecules.
- (ii) Uses:
  - As fuels, as starting materials for synthesis. Uses of haloalkanes and pollution effects.
(g) Alkenes:
- (i) Sources and properties:
  - Laboratory preparation.
  - Nomenclature and structure.
  - Addition reactions with halogens hydrogen, bromine water, hydrogen halides, and acidified water.
  - Oxidation: hydroxylation with aqueous KMnO_4.
  - Polymerization.
- (ii) Uses:
- (iii) Laboratory detection:
  - Use of reaction with Br2/water, Br2/CCl4 and KMnO4(aq) as means of characterizing alkenes.
(h) Alkynes:
- (i) Sources, characteristic properties, and uses:
  - Nomenclature and structure.
  - Industrial production of ethyne.
  - Uses of ethyne.
  - Distinguishing test between terminal and non-terminal alkynes.
  - Test to distinguish between alkane, alkene, and alkyne.
- (ii) Chemical reactions:
  - Halogenation, combustion, hydration, and hydrogenation.
(i) Benzene:
- Resonance in benzene. Stability leading to substitution reactions.
  - (i) Structure and physical properties:
  - (ii) Chemical properties:
    - Addition reactions: hydrogenation and halogenation (mechanism not required).
    - Compare reactions with those of alkenes.
(J) Alkanols:
- (i) Sources, nomenclature, and structure:
  - Laboratory preparation, including hydration of alkenes.
  - Industrial and local production of ethanol, including alcoholic beverages,
  - Harmful impurities and methods of purification should be mentioned.
  - Recognition of the structure of mono-, di-, and triols.
- (ii) Classification:
  - Primary, secondary, and tertiary alkanols.
- (iii) Physical properties:
  - Boiling point, solubility in water. Including hydrogen bonding effect.
- (iv) Chemical properties:
  - Reaction with:
    - Na;
    - alkanoic acids (esterification);
    - conc. H2SO4.
  - Oxidation by:
    - KMnO_4(aq);
    - K2Cr2O_7(aq);
    - I_2 in NaOH(aq).
- (v) Laboratory test:
  - Laboratory test for ethanol.
- **(vi) Uses.
(k) Alkanoic acids:
- (i) Sources, nomenclature, and structure:
  - Methanoic acid – insect bite.
  - Ethanoic acid – vinegar.
  - Recognition of mono and dioic acid.
- (ii) Physical properties:
  - Boiling point, solubility in water. Including hydrogen bonding effect.
- (iii) Chemical properties:
  - Acid properties only i.e., reactions with H2O, NaOH, NH3, NaHCO_3, Zn, and Mg.
- **(iv) Laboratory test:
- (iv) Uses:**
  - Reaction with NaHCO3, Na2CO_3.
  - Uses of ethanoic and phenyl methanoic (benzoic) acids as examples of aliphatic and aromatic acids, respectively.
(l) Alkanoates as derivatives of alkanoic acids:
- (i) Sources, nomenclature, preparation, and structure:
  - Preparation of alkyl alkanoates (esters) from alkanoic acids.
- (ii) Physical properties:
  - Solubility, boiling, and melting point.
- (iii) Chemical properties:
  - Hydrolysis of alkyl alkanoates (mechanism not required).
- (iv) Uses:
  - Uses of alkanoates to include production of soap, flavoring agent, plasticizers, as solvents, and in perfumes.

14.0 CHEMISTRY, INDUSTRY, AND THE ENVIRONMENT

(a) Chemical industry:
- Natural resources in the candidate’s own country.
- Chemical industries in candidates own country and their corresponding raw materials.
- Distinction between fine and heavy chemicals.
- Factors that determine the location of chemical industries.
- The effect of industries on the community.
(b) Pollution: air, water and soil pollution:
- Sources, effects, and control.
- Greenhouse effect and depletion of the ozone layer.
- Biodegradable and non-biodegradable pollutants.
(c) Biotechnology:
- Food processing, fermentation including production of gari, bread, and alcoholic beverages e.g., Local gin.

15.0 BASIC BIOCHEMISTRY AND SYNTHETIC POLYMERS

(a) Proteins:
- Proteins as polymers of amino acids molecules linked by peptide or amide linkage.
  - (i) Sources and properties:
    - Physical properties e.g., solubility.
  - (ii) Uses of protein:
    - Chemical properties to include:
      - hydrolysis of proteins;
      - laboratory test using Ninhydrin/Biuret reagent/Millons reagent.
(b) Amino acids:
- Nomenclature and general structure of amino acids.
- Difunctional nature of amino acids.
(c) Fats/oils:
- As alkyl alkanoates (esters).
  - (i) Sources and properties:
    - From animals and plants.
    - Physical properties such as solubility.
  - (ii) General structure of fats/oils:
    - Chemical properties:
      - acidic and alkaline hydrolysis;
      - hydrogenation;
      - test for fats and oil.
    - As mono-, di-, and tri- esters of propane-1,2,3-triol (glycerol).
  - (iii) Preparation of soap:
    - Preparation of soap (saponification) from fats and oils.
    - Comparison of soap less detergents and their action on soft and hard water.
  - **(iv) Uses of fats/oils.
(d) Carbohydrates:
- (i) Sources and nomenclature:
  - Classes of carbohydrates as:
    - monosaccharides;
    - disaccharides;
    - polysaccharides.
  - Name and components of various classes of carbohydrates.
- (ii) Properties: